{"id":238,"date":"2018-11-21T18:36:20","date_gmt":"2018-11-21T18:36:20","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/suny-potsdam-organicchemistry2\/?post_type=chapter&p=238"},"modified":"2019-01-08T14:38:09","modified_gmt":"2019-01-08T14:38:09","slug":"13-4-aromaticity","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/suny-potsdam-organicchemistry2\/chapter\/13-4-aromaticity\/","title":{"raw":"13.6. Aromaticity","rendered":"13.6. Aromaticity"},"content":{"raw":"
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Aromaticity is a property of conjugated cycloalkenes in which the stabilization of the molecule is enhanced due to the ability of the electrons in the $$\\pi$$ orbitals to delocalize. This act as a framework to create a planar molecule.\r\n
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Introduction<\/h3>\r\nWhy do we care if a compound is aromatic or not? Because we encounter aromatics every single day of our lives. Without aromatic compounds, we would not only be lacking many material necessities, our bodies would also not be able to function. Aromatic compounds are essential in industry; about 35 million tons of aromatic compounds are produced in the world every year to produce important chemicals and polymers, such as polyester and nylon. Aromatic compounds are also vital to the biochemistry of all living things. Three of the twenty amino acids<\/a> used to form proteins (\"the building blocks of life\") are aromatic compounds and all five of the nucleotides that make up DNA <\/a>and RNA sequences are all aromatic compounds. Needless to say, aromatic compounds are vital to us in many aspects.\r\n
The three general requirements for a compound to be aromatic are:<\/div>\r\n
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  1. The compound must be cyclic<\/li>\r\n \t
  2. Each element within the ring must have a p-orbital that is perpendicular <\/em>to the ring, hence the molecule is planar.\u00a0 This means that the ring cannot contain a neutral sp3<\/sup> carbon.<\/li>\r\n \t
  3. The compound must follow H\u00fcckel's Rule<\/a> (the ring has to contain 4n+2 p-orbital electrons).<\/li>\r\n<\/ol>\r\nAmong the many distinctive features of benzene<\/a>, its aromaticity is the major contributor to why it is so unreactive. This section will try to clarify the theory of aromaticity and why aromaticity gives unique qualities that make these conjugated alkenes inert to compounds such as Br2<\/sub> and even hydrochloric acid. It will also go into detail about the unusually large resonance energy due to the six conjugated carbons of\u00a0benzene.\r\n\r\n\"dd.bmp\"\r\n\r\nThe delocalization of the p-orbital carbons on the sp2\u00a0<\/sup>hybridized carbons is what gives the aromatic qualities of benzene.\r\n\r\n<\/div>\r\n<\/div>\r\n
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    Basic Structure of Benzene<\/h3>\r\n\"Benzene\r\n\r\nBecause of the aromaticity of benzene, the resulting molecule is planar in shape with each C-C bond being 1.39 \u00c5 in length and each bond angle being 120\u00b0. You might ask yourselves how it's possible to have all of the bonds to be the same length if the ring is conjugated with both single (1.47 \u00c5) and double (1.34 \u00c5), but it is important to note that there are no distinct single or double bonds within the benzene. Rather, the delocalization of the ring makes each count as one and a half bonds between the carbons which makes sense because experimentally we find that the actual bond length is somewhere in between a single and double bond. Finally, there are a total of six p-orbital electrons that form the stabilizing electron clouds above and below the aromatic ring.\r\n\r\n<\/div>\r\n<\/div>\r\n
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    Evidence of aromaticity: Heats of hydrogenation<\/h3>\r\nOne of the ways to test the relative amounts of resonance energy in a molecule is to compare the heats of hydrogenation between similar compounds. For instance, if we compare cyclohexene, 1,3-cyclohexadiene, and benzene, we would expect that their heats of hydrogenation will increase since the number of double bonds increases respectively. However, experimental evidence suggests that the actual heat of hydrogenation for benzene is actually 49.3 kcal\/mole, making it even more stable than the 1,3-cyclohexadiene even though it has two double bonds, compared to benzene's three double bonds. This characteristic can be attributed to the aromaticity of benzene which delocalizes the electrons of the six pi orbitals.\r\n\r\n\"cc\r\n\r\n<\/div>\r\n<\/div>\r\n
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    References<\/h3>\r\n
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    1. P. Schleyer, \"Aromaticity (Editorial)\", Chemical Reviews, 2001.<\/li>\r\n \t
    2. P. Schleyer, \"Introduction: Delocalization-? and ? (Editorial)\", Chemical Reviews,\u00a0 2005.<\/li>\r\n \t
    3. Vollhardt, K. Peter C., and Neil E. Schore. Organic Chemistry: Structure and Function. New York: W.H. Freeman and Company, 2007.<\/li>\r\n<\/ol>\r\n<\/div>\r\n<\/div>\r\n
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      Exercises<\/h3>\r\n
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      1. What is the hybridization of each carbon and the overall shape of benzene?<\/li>\r\n \t
      2. What is the resonance energy of benzene?<\/li>\r\n \t
      3. Place the following compounds in order of heats of hydrogenation from smallest to greatest : Benzene, 1,3-Cyclohexadiene, and Cyclohexene.<\/li>\r\n<\/ol>\r\n<\/div>\r\n<\/div>\r\n
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        Answers<\/h3>\r\n[reveal-answer q=\"378415\"]Show Answer[\/reveal-answer]\r\n[hidden-answer a=\"378415\"]\r\n
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        1. All six carbons are sp2\u00a0 hybridized and the aromaticity of the benzene creates a planar molecule.<\/li>\r\n \t
        2. 29.6 kcal\/mol<\/li>\r\n \t
        3. Cyclohexene < Benzene < 1,3-Cyclohexadiene.[\/hidden-answer]<\/li>\r\n<\/ol>\r\n<\/div>\r\n<\/div>\r\n
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          What makes a structure aromatic?<\/h2>\r\nValence bond theory uses Lewis diagrams to depict structure and bonding of covalent entities, such as molecules and polyatomic ions, henceforth referred to just as molecules. The Lewis diagram of many a molecule, however, is not consistent with the observed properties of the molecule. \u00a0The Lewis diagram of some molecules suggests a ring bearing a fully conjugated $$\\pi$$-electron system, or loop of $$\\pi$$ electrons, provided each atom in the ring is sp2<\/sup>- or sp-hybridized.\u00a0 Note that in cases like 5 that contain sp3 atoms with lone pairs, these atoms (usually N, O or S) can rehybridize to sp2<\/sup> and make the remaining p-orbital available for bonding.\u00a0 A fully bonded sp3<\/sup> carbon cannot do this, and thus cannot form part of a fully conjugated system.\r\n